Question

If standard reduction potential ($E^\circ$) of ($Mg^{2+} | Mg(s)$), ($Ag^+ | Ag(s)$), ($Zn^{2+}(aq) | Zn(s)$) and ($Cu^{2+}(aq) | Cu(s)$) are $-2.37$ V, $+0.79$ V, $-0.76$ V and $+0.34$ V respectively. Which of the following reaction is spontaneous?

• A. $Zn(s) + Mg^{2+}(aq) \rightarrow Zn^{2+}(aq) + Mg(s)$
• B. $2Ag(s) + Zn^{2+}(aq) \rightarrow 2Ag^+(aq) + Zn(s)$
• C. $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$
• D. $Cu(s) + Mg^{2+}(aq) \rightarrow Cu^{2+}(aq) + Mg(s)$

Hint:

The Electrochemical Series rule of thumb: "The lower the $E^\circ$ value, the stronger the reducing agent." A metal can only displace another metal if it sits below it in the standard reduction potential table.

Answer 1

The Correct Option is C

Step 1: Understanding the Concept:
A redox reaction is spontaneous if the cell potential ($E_{cell}^\circ$) is positive. $E_{cell}^\circ = E_{cathode}^\circ - E_{anode}^\circ$.

Step 2: Formula Application:
The metal with the higher (more positive) $E^\circ$ acts as the cathode (reduction), and the metal with the lower (more negative) $E^\circ$ acts as the anode (oxidation).

Step 3: Explanation:
For option (c): $Zn$ is oxidized (anode) and $Cu$ is reduced (cathode). $E_{cell}^\circ = E_{Cu}^\circ - E_{Zn}^\circ = (+0.34) - (-0.76) = +1.10$ V. Since $E_{cell}^\circ > 0$, the reaction is spontaneous. In all other options, the metal with the higher reduction potential is being oxidized, leading to a negative $E_{cell}^\circ$.

Step 4: Final Answer:
The spontaneous reaction is (c).